Equilibrium

The Equilibrium Constant, K_EQ

    In equilibrium, an important concept is that of the reversible reaction, one whose reactants form products, but also can have the products form the reactants.  A reaction is said to be at equilibrium when the rate of the foward reaction is equal to the rate of the reverse reaction.  The folowing is the general expression for the equilibrium constant:

       aA + bB <---> cC + dD

                    [C]^c[D]^d
        K_eq=   ---------
                    [A]^a[B]^b

        1.  The products are in the numerator, reactants in the denominatior
        2.  Coeffficients of balanced equation become exponents
        3.  Solids and pure Liquids are ignored
        4.  There are no units assigned to K_eq

K_eq and Gasses

    Previously we have shown you how to calculate K_eq in terms of concentration (or K_c, as we will refer to it).  Now if you would like to use gasses, you might want to express K_eq in terms of partial pressures (or K_p, as we will refer to it).  Here is an equation that you will be able to use.

          K_p  = K_c(RT)ⁿ

        R = the ideal gas constant, 0.0821 (L-atm)/(mol-K)
        T = absolute temperature (K)
        n = change in moles (Product - Reactant)

The Reaction Quotient, Q

    The reaction quotient is calculated exactly the same as the equilibrium constant, except that the inital conditions are used instead of the conditions at equilibrium.  The purpose of doing this is to test which way the reaction is going to go.  If Q is less than the calculated K for a reaction, then the reaction will proceed towards fowared, generating more products.  If Q is greater than K, then the reaction will go backwards, generating more reactants.  And finally, if Q = K than the reaction is at equilibrium. 

K_eq and Multistep Processes

    The simple rule for this is that if there are two reactions that can be added togeather, then to find the K_eq final for the reaction you just multiply.  Here is an example:

            A <--> B + C             K_eq = K₁
            B + D <--> E             K_eq = K₂
            A + C + D <--> E      K_eq = K₁K₂

Le Chatelier's Principle

    This principle deals with added stresses of reactions, and the shifting that occurs because of them. There are three things that can cause a shift in equilibrium because of stresses.

Concentration Change

    When the concentration of a reactant or product is increased, then the reaction will then go towards the other side of the equation.  For instance, if a reactant is added to the solution, then the equilibrium will then shift towards the products.  This occurs so the reaction can then use up the extra or added substance.

    When the concentration of a reactant or product is decreased, then the reaction will then go towards the same side of the equation.  For instance, if a product is taken out of solution then the equilibrium will shift towards the products.   This will happen to replace the lost substance.

Volume Change

    When the volume is increased in a solution, the equilibrium will shift towards the side that produces more moles of gas.  Usually this will be the side where the coefficents add up to a higher number.  The oppsite is also true.

    If there is no gas involved in the given equation, or the number of moles are the same on both sides of the equation, than changing the volume will have no effect on shifting the equilibrium.

Temperature Change

    When the temperature is increased, the reaction will procede towards the endothermic direction.  The opposite is also true.  For instance, if temperature is increased and the delta H is negitive, the reaction will procede towards the reactants.  In that same case if the delta H was positive, then it would go towards the products.